Coordination compounds, such as the FeCl4-
ion and CrCl3 6 NH3, are called such
because they contain ions or molecules linked, or coordinated, to
a transition metal. They are also known as complex ions or
coordination complexes because they are Lewis acid-base
complexes. The ions or molecules that bind to transition-metal
ions to form these complexes are called ligands (from
Latin, "to tie or bind"). The number of ligands bound
to the transition metal ion is called the coordination number.

Although coordination complexes are particularly important in
the chemistry of the transition metals, some main group elements
also form complexes. Aluminum, tin, and lead, for example, form
complexes such as the AlF63-, SnCl42-
and PbI42- ions.

Alfred Werner developed a model of coordination complexs which
explains the following observations.

At least three different cobalt(III) complexes can be
isolated when CoCl2 is dissolved in aqueous
ammonia and then oxidized by air to the +3 oxidation
state. A fourth complex can be made by slightly different
techniques. These complexes have different colors and
different empirical formulas.

CoCl3 6 NH3

orange-yellow

CoCl3 5 NH3 H2O

red

CoCl3 5 NH3

purple

CoCl3 4 NH3

green

The reactivity of the ammonia in these complexes has been
drastically reduced. By itself, ammonia reacts rapidly
with hydrochloric acid to form ammonium chloride.

NH3(aq) +
HCl(aq) NH4+(aq) + Cl-(aq)

These complexes don't react with hydrochloric acid, even at
100oC.

CoCl3 6 NH3(aq)
+ HCl(aq)

Solutions of the Cl- ion react with Ag+
ion to form a white precipitate of AgCl.

Ag+(aq) +
Cl-(aq) AgCl(s)

When excess Ag+ ion is added to solutions of the
CoCl3 6 NH3 and CoCl3 5 NH3
H2O complexes, three moles of AgCl are formed for
each mole of complex in solution, as might be expected. However,
only two of the Cl- ions in the CoCl3 5
NH3 complex and only one of the Cl- ions in
CoCl3 4 NH3 can be precipitated with Ag+
ions.

Measurements of the conductivity of aqueous solutions of
these complexes suggest that the CoCl3 6 NH3
and CoCl3 5 NH3 H2O
complexes dissociate in water to give a total of four
ions. CoCl3 5 NH3 dissociates to
give three ions, and CoCl3 4 NH3
dissociates to give only two ions.

Werner explained these observations by suggesting that
transition-metal ions such as the Co3+ ion have a
primary valence and a secondary valence. The primary valence
is the number of negative ions needed to satisfy the charge on
the metal ion. In each of the cobalt(III) complexes previously
described, three Cl- ions are needed to satisfy the
primary valence of the Co3+ ion.

The secondary valence is the number of ions of
molecules that are coordinated to the metal ion. Werner assumed
that the secondary valence of the transition metal in these
cobalt(III) complexes is six. The formulas of these compounds can
therefore be written as follows.

[Co(NH3)63+][Cl-]3

orange-yellow

[Co(NH3)5(H2O)3+][Cl-]3

red

[Co(NH3)5Cl2+][Cl-]2

purple

[Co(NH3)4Cl2+][Cl-]

green

The cobalt ion is coordinated to a total of six ligands in
each complex, which satisfies the secondary valence of this ion.
Each complex also has a total of three chloride ions that satisfy
the primary valence. Some of the Cl- ions are free to
dissociate when the complex dissolves in water. Others are bound
to the Co3+ ion and neither dissociate nor react with
Ag+.

The [Co(NH3)6]Cl3 complex
dissociates in water to give a total of four ions, and all three
Cl- ions are free to react with Ag+ ion.

H2O

[Co(NH3)6]Cl3(s)

Co(NH3)63+(aq)
+ 3 Cl-(aq)

One of the chloride ions is bound to the cobalt in the [Co(NH3)5Cl]Cl2
complex. Only three ions are formed when this compound dissolves
in water, and only two Cl- ions are free to
precipitate with Ag+ ions.

H2O

[Co(NH3)5Cl][Cl]2(s)

Co(NH3)5Cl2+(aq)
+ 2 Cl-(aq)

Once again, the three Cl- ions are free to
dissociate when [Co(NH3)5(H2O)]Cl3
dissolves in water, and they precipitate when Ag+ ions
are added to the solution.

H2O

[Co(NH3)5(H2O)]Cl3(s)

Co(NH3)5(H2O)3+(aq)
+ 3 Cl-(aq)

Two of the chloride ions are bound to the cobalt in [Co(NH3)4Cl2]Cl.
Only two ions are formed when this compound dissolves in water,
and only one Cl- ion is free to precipitate with Ag+
ions.

H2O

[Co(NH3)4Cl2][Cl](s)

Co(NH3)4Cl2+(aq)
+ Cl-(aq)

Werner assumed that transition-metal complexes had definite
shapes. According to his theory, the ligands in six-coordinate
cobalt(III) complexes are oriented toward the corners of an
octahedron, as shown in the figure below.

Any ion or molecule with a pair of nonbonding electrons can be
a ligand. Many ligands are described as monodentate
(literally, "one-toothed") because they
"bite" the metal in only one place. Typical monodentate
ligands are given in the figure below.

Other ligands can attach to the metal more than once.
Ethylenediamine (en) is a typical bidentate ligand.

Each end of this molecule contains a pair of nonbonding
electrons that can form a covalent bond to a metal ion.
Ethylenediamine is also an example of a chelating ligand.
The term chelate comes from a Greek stem meaning
"claw." It is used to describe ligands that can grab
the metal in two or more places, the way a claw would.

Linking ethylene- diamine fragments gives tridentate
ligands and tetradentate ligands, such as
diethylenetriamine (dien) and triethylenetetramine (trien).
Adding four -CH2CO2- groups to
an ethylenediamine framework gives a hexadentate ligand,
which can single-handedly satisfy the secondary valence of a
transition-metal ion.

Transition-metal complexes have been characterized with
coordination numbers that range from 1 to 12, but the most common
coordination numbers are 2, 4, and 6. Examples of complexes with
these coordination numbers are given in the table below.

Examples of Common Coordination Numbers

Metal Ion

Ligand

Complex

CoordinationNumber

Ag+

+

2 NH3

Ag(NH3)2+

2

Ag+

+

2 S2O32-

AgCl2-

2

Ag+

+

2 Cl-

Ag(S2O3)23-

2

Pb2+

+

2 OAc-

Pb(OAc)2

2

Cu+

+

2 NH3

Cu(NH3)2+

2

Cu2+

+

4 NH3

Cu(NH3)42+

4

Zn2+

+

4 CN-

Zn(CN)42-

4

Hg2+

+

4 I-

HgI42-

4

Co2+

+

4 SCN-

Co(SCN)42-

4

Fe2+

+

6 H2O

Fe(H2O)62+

6

Fe3+

+

6 H2O

Fe(H2O)63+

6

Fe2+

+

6 CN-

Fe(CN)64-

6

Co3+

+

6 NH3

Co(NH3)63+

6

Ni2+

+

6 NH3

Ni(NH3)62+

6

Note that the charge on the complex is always the sum of the
charges on the ions or molecules that form the complex.

Cu2+ + 4 NH3 Cu(NH3)42+

Pb2+ + 2 OAc- Pb(OAc)2

Fe2+ + 6 CN- Fe(CN)64-

Note also that the coordination number of
a complex often increases as the charge on the metal ion becomes
larger.

Cu+ + 2 NH3 Cu(NH3)2+

Cu2+ + 4 NH3
Cu(NH3)42+

Practice Problem 2:

Calculate
the charge on the transition-metal ion in the following
complexes.

G. N. Lewis was the first to recognize that the reaction
between a transition-metal ion and ligands to form a coordination
complex was analogous to the reaction between the H+
and OH- ions to form water. The reaction between H+
and OH- ions involves the donation of a pair of
electrons from the OH- ion to the H+ ion to
form a covalent bond.

The H+ ion can be described as an electron-pair
acceptor. The OH- ion, on the other hand, is an electron-pair
donor. Lewis argued that any ion or molecule that behaves
like the H+ ion should be an acid. Conversely, any ion
or molecule that behaves like the OH- ion should be a
base. A Lewis acid is therefore any ion or molecule that
can accept a pair of electrons. A Lewis base is an ion or
molecule that can donate a pair of electrons.

When Co3+ ions react with ammonia, the Co3+
ion accepts pairs of nonbonding electrons from six NH3
ligands to form covalent cobalt-nitrogen bonds as shown in the
figure below.

The metal ion is therefore a Lewis acid, and the ligands
coordinated to this metal ion are Lewis bases.

The Co3+ ion is an electron-pair acceptor, or Lewis
acid, because it has empty valence-shell orbitals that can be
used to hold pairs of electrons. To emphasize these empty valence
orbitals we can write the configuration of the Co3+
ion as follows.

Co3+: [Ar] 3d6 4s0
4p0

There is room in the valence shell of this ion for 12 more
electrons. (Four electrons can be added to the 3d
subshell, two to the 4s orbital, and six to the 4p
subshell.) The NH3 molecule is an electron-pair donor,
or Lewis base, because it has a pair of nonbonding electrons on
the nitrogen atom.

According to this model, transition-metal ions form
coordination complexes because they have empty valence-shell
orbitals that can accept pairs of electrons from a Lewis base.
Ligands must therefore be Lewis bases: They must contain at least
one pair of nonbonding electrons that can be donated to a metal
ion.